As we read across the first-row transition metals from left to right, the +2 oxidation state becomes more stable in comparison with the +3 state. Why is this so?

The starting point to grasping the behavior of transition metals lies in their electron configuration. For transition elements, the d-orbitals – the set of five orbitals with similar energy – become progressively filled with electrons as you move from left to right across the periodic table. At the outset, these elements have their outermost s-orbital filled with two electrons. For instance, scandium, the first transition metal, has an electron configuration of [Ar] 4s^2 3d^1.

When an element loses two electrons to form a +2 oxidation state, it often loses the s-orbital electrons first, leading to a configuration featuring an increasingly stable d-orbital electron arrangement. As you continue across the period, additional electrons fill the d-orbitals, which influences the chemistry and stability of the metals' different oxidation states. The +2 state remains favorable because it avoids destabilizing the d-orbital electron arrangement. This is a key concept in understanding why +2 oxidation state is more prevalent among the heavier transition metals compared to the lighter ones with fewer d-electrons.

The stability of an electron configuration, particularly in d-orbitals, is central to a transition metal's chemical properties. In the context of d-orbital stability, 'half-filled' (where there are five electrons) and 'fully-filled' (where there are ten electrons) configurations are particularly stable. This is due to the symmetrical distribution of electrons and the exchange energy that stabilizes the configuration.

For example, a metal like manganese, with an electron configuration of [Ar] 4s^2 3d^5,

has a particularly stable +2 oxidation state as it corresponds to a half-filled d-orbital. When a metal ion transitions to a +3 state, it often requires removing an electron from a more stable d-orbital arrangement, leading to increased instability and higher energy costs. This factor is why as you progress towards the right in the periodic table, metals with more d-electrons will favor the +2 oxidation state to preserve d-orbital stability.

Oxidation state trends for transition metals are not arbitrary; they are closely linked to electronic structure and d-orbital stability. As we move from left to right across the transition series, the +2 oxidation state typically becomes more stable relative to the +3 state. This trend arises because removing a third electron after the +2 state often involves taking an electron from an increasingly stable d-orbital.

In elements towards the left of the series, there aren't as many d-electrons to remove; hence, achieving a +3 or even higher oxidation states doesn't result in significant destabilization. However, for elements on the right, having more d-electrons means that achieving a +3 state would require removing electron(s) that contribute greatly to the current stability of the d-orbital electron cloud. This process requires additional energy, making the +3 state less favorable. Consequently, as we progress across the period, the predominance of the +2 oxidation state over the +3 state becomes one of the quintessential trends in transition metal chemistry.